1. Electrons in Atoms
Chapter 4

2. The New Atomic Model
Investigations  relationship between light and atom’s electrons
How are electrons arranged? Why don’t they fall into the nucleus?

3. Light a wave or particle?
Wave Description:
Electromagnetic Radiation: energy that acts like a wave in space
All forms create Electromagnetic Spectrum

4. Electromagnetic Spectrum

6. Electromagnetic Spectrum
All waves move at speed of light, c, 3.00x108 m/s
Waves identified by:
wavelength, , the distance b/ corresponding points on adjacent waves. Units: nm, cm, or m
Frequency, , # of waves that pass a given point in a specific time, 1 sec. Unit: 1/s = Hertz, Hz

7. Wavelength and Frequency

8. Wavelength and Frequency
Inverse proportion equation!!
Frequency, 1/s
speed of light, m/s
wavelength, m

9. Calculation
Calculate the wavelength of a radio wave with a frequency of x 106s-1
Determine the frequency of light whose wavelength is nm.

10. Particle Nature of Light
Photoelectric Effect: emission of electrons from a metal when light shines on the metal

11. Photoelectric Effect
Light had to be certain frequency to knock e- loose
Light must also be a particle!
Max Planck(1900) explanation: objects emit energy in small packets called quanta.
A photon is a single quantum of (visible) light as well as a single quantum of all other forms of electromagnetic radiation, and can be referred to as a "light quantum".

12. Max Planck
Quantum of energy is the smallest amount of energy that can be lost or gained by an atom
E = h
Frequency, s-1
Energy of quantum, in joules, J
Planck’s constant, 6.626x10-34 Js

13. Energy Calculation
What is the energy of green light, with a wavelength of 500. nm?
500. nm = 5.00 x 10-7 m
1 nm = 10-9 m

14. Electromagnetic Spectrum
What is the color of an electromagnetic radiation with an energy of 3.4 x10-19 J?

15. Photoelectric Effect Albert Einstein Light is both wave and particle!
Particle of light = photon, having zero mass and a quantum of energy
Photons hit metal and knock e- out, but photon has to have enough energy

16. H-atom Emission Spectrum
Pass a current through gas at low pressure it excites the atoms
Ground state: lowest energy state of an atom
Excited state: atom has higher potential energy than it has in ground state

17. H – Atom Spectrum
When atom jumps from excited state to ground state it gives off energy  LIGHT! E2
Ephoton = E2 – E1 = hv E1

18. Bohr’s Model of the H Atom (and only H!)
Applied quantization of energy transfer to the atomic model
Studied atomic spectrum of H to come up with atomic model.
Atomic emission spectra:
Most sources produce light that contains many wavelengths at once.
However, light emitted from pure substances may contain only a few specific wavelengths of light called a line spectrum (as opposed to a continuous spectrum).
Atomic emission spectra are inverses of atomic absorption spectra.

19. Atomic Emission Spectra of C and H
Hydrogen: contains 1 red, 1 green, 1 blue and 1 violet.
Carbon: Contains many more emission lines as compared to H. Why?

20. H-atom Line Emission Spectrum

21. Element Emission Spectras
Helium – 23 lines
Neon – 75 lines
Argon lines
Xenon – 139 lines
Mercury – 40 lines

22. H-atom Line Emission Spectrum
More lines in UV (Lyman series) and IR(Paschen series)
Why did H-atom only emit certain colors of light?

23. Bohr Model of H-atom 1913 – Niels Bohr
e- circles nucleus in certain paths, orbits or atomic energy levels
e- is higher in energy the farther away from nucleus
e- cannot be between orbits
Video - 23

24. Bohr Model of H-atom

25. Bohr Model of H-atom
From wavelengths of emission spectrum Bohr calculated energy levels of H-atom
Model worked ONLY for H-atom
End Part 1